Samples Containing Iron

General Information

Iron is the fourth most abundant element in the earth's crust, after oxygen, silicon, and aluminum. Iron has been known by man back to prehistoric times, and was used for making weapons and tools. During the Middle Ages, high-grade steel was manufactured in Arabia and Spain. Although pure Fe finds little commercial use, the addition of a few elements makes it arguably the most important material used in industry. Fe is also one of, if not the most, common elements requiring analysis by analytical laboratories that have a variety of sample preparation and measurement techniques available.

Iron is found in the +2 and +3 oxidation states, and is commonly found as iron (II) oxide (FeO); iron (III) oxide (Fe₂O₃); magnetite (Fe₃O₄); limonite (FeO(OH)·nH₂O); siderite (FeCO₃); pyrrhotite (Fe_1-xS); and a large number of combinations with many other elements, such as chromite (FeO·Cr₂O₃). The chemical resistance to bringing about a solution varies widely, with the freshly precipitated oxides/hydroxides dissolving with little effort, and ignited oxides and combinations like the chromite requiring a much greater effort. The techniques proposed here are applicable to this wide spectrum of sample types.

In general, Fe compounds are not easy to dissolve, and their chemical behavior is not that predictable. Fe exists in oxidation states ranging from +6, +5, +4, +3, +2, +1, and 0, with the +3, +2, and 0 states being by far the most common. Consequently, the chemistry of iron addressed in this chapter will be limited to these three oxidation states. The literature addressing the analytical chemistry of Fe is extensive. This discussion will present methods that have been performed for intended analysis by ICP-OES, ICP-MS, and AAS, but that should also be applicable to Fe analysis using colorimetric, titrimetric, gravimetric, or electrochemical methods of analysis. Questions concerning the applicability of a given preparation technique to an intended measurement method are welcome.

In acidic media, Fe³⁺ is thermodynamically favored, and is the state that results from most sample preparations intended for ICP measurement. In some preparations, Fe will go into solution as Fe²⁺, but will slowly be oxidized by dissolved oxygen to Fe³⁺, unless the solution is purged and blanketed with high-purity nitrogen. For example, this is the case when Fe metal is dissolved with HCl. The following reduction potentials are very helpful in explaining the chemistry of Fe and its oxidation states (taken from Electroanalytical Chemistry, 2nd edition, James J. Lingane, Interscience Publishers, NY, 1970):

The +2 oxidation state of Fe is much more stable in acidic than in basic media to oxidation to the +3 state. Standard solutions of Fe²⁺ can be prepared, but are more stable if air is excluded. When air exclusion is not possible, then they should be restandardized each analytical day. Fe²⁺ hydrolyzes only slightly before precipitation of the Fe(OH)₂, and this precipitation does not begin until a pH of ≥6.5. Fe²⁺ can be formed with non-oxidative sample preparations, but with slow air oxidation to Fe³⁺. Fe³⁺ begins precipitation as the hydroxide at a pH of ~0.5. Unlike the +2 state, hydrolysis of Fe³⁺ ions involves the formation of polynuclear complexes, such as Fe₃(OH)₄⁵⁺. The form in which the Fe³⁺ hydroxide precipitates depends on conditions such as the pH, temperature, and anions present in solution, and is evidenced by a wide variety of shades of brown to orange. The formation of complexes with both the +2 and +3 oxidation states results in an ability to maintain stable solutions with increasing pH, and also results in a lowering of the reduction potentials. For example, a 0.1M Fe/EDTA complex is stable at pH values of 11.5, and that of 0.1M cyanide is stable at a pH >12, which is rather remarkable considering the insolubility of Fe(OH)₃. When Fe(OH)₂ is precipitated in the presence of air, it turns black immediately, due to the formation of the +3 oxidation state and a magnetite-like, mixed-oxidation-state material.

Salts

Fe(II) salts, in crystals and in solution, have a pale green color. Solutions of the salts are slightly acidic. Fe(III) salts form solutions with a brownish-yellow color. The Fe(II) sulfate heptahydrate (48), chloride (62), bromide (117), and nitrate hexahydrate (77) are readily water-soluble (solubility in g/100 g water shown in parentheses), with the Fe(III) salts of these same anions showing similar solubility. The carbonate (Fe²⁺), oxides, and sulfides of both oxidation states are insoluble in water, and the fluorides are only slightly water-soluble. Solutions of Fe³⁺, when treated with alkali hydroxides, all form Fe(OH)₃, which is a reddish-brown solid that does not redissolve in an excess of the hydroxide, unlike its neighbor Cr, or Co, Ni, and Zn, which will dissolve in excess NH₄OH. Fe(OH)₃ is used in trace analysis as a gathering agent, because its gelatinous property and structure bring about co-precipitation of trace metals in the same solution.

Sampling and Handling

The measurement of Fe is required for a wide range of samples, including ores and minerals, soils, sludges, water (both drinking and waste), biological, agricultural, metallurgical, and industrial samples. There is a great risk of contamination when any sample handling apparatus/device is made of, or has come in contact with, stainless steel. In addition, since Fe is so prevalent in the earth's crust, the air (particulate dust) must be considered as a contamination route. The following considerations apply:

• Many tools that pulverize, mix, cut, etc., contain Fe. Attempt to use devices made of ceramics, silica/quartz, and polymers where possible.
• Many polymers, such as Teflon, have been exposed to stainless steel (cutting/pulverizing) prior to molding. New plastic and Teflon containers should always be acid-leached to remove possible contamination from the stainless steel elements.
• The collection of biological samples is also at risk of contamination, due to the very low (ppb) levels of Fe sought. The use of steel needles and scalpels, or any metallic object that may contain Fe, should be avoided.

The risk of contamination is great for Fe when alloys, steels, and grinding equipment are used in some part of the sample collection or preparation. For more on sample contamination risks, see chapters 8, 9, and 10 of the Inorganic Ventures Trace Analysis Guide.

For general information on sampling and subsampling, see Sampling and Subsampling.

The Metal and Alloys

Metal — Iron metal dissolves in HCl and dilute H₂SO₄, forming Fe²⁺ and liberation of H₂. See the following equation:

Fe⁰ + H₂SO₄ = FeSO₄ + H₂

Concentrated cold sulfuric acid has no action, but if hot, SO₂ is evolved and Fe³⁺ formed, as per the following equation:

2 Fe⁰ + 6 H₂SO₄ = Fe₂(SO₄)₃ + 3 SO₂ + 6 H₂O

Fe metal will not dissolve in cold, concentrated nitric acid, but when heated to near boiling, it dissolves vigorously with the evolution of brown NO₂.

Alloys — The use of nitric + HCl is most commonly used for Cr/Ni and conventional steels, and the addition of water to this mixture is helpful in increasing the sample size and will allow for the preparation of the Fe/Ni high-temperature alloys, as well as conventional steels. When an element is present in the steel/alloy, such as Si, Ti, Ta, Zr, W, etc., the addition of HF is essential, and the presence of HCl, although often used, becomes less critical. A typical preparation for a Zr steel (0.1 grams of alloy/steel + 20 mL conc. HNO₃ + 10 mL 1:1 HCl + 5 mL HF) is somewhat different than that needed for a ferrosilicon (0.1 gram sample + 20 mL conc. HNO₃ + 5 mL 40% HF), with the understanding that just about any alloy of Fe will dissolve in a mixture of nitric/HCl/HF, with the addition of water generally increasing the amount of sample that can be dissolved, but with too much water quenching the entire reaction. An impressive mixture is a 1:1:1 mixture of HNO₃ + HF + H₂O, used to dissolve high-temperature alloys of Fe/Co/Cr/Ti/Al/Mo/Ta/Hf (use 10 mL mixture/0.2 grams sample).

Oxides, Minerals and Ores

Oxides and hydroxides — Fe₂O₃, after ignition to temperatures at or above typical ashing temperatures (450-1000 °C), is rendered insoluble, or at best very slightly soluble, in acids. HCl is the best solvent, but dissolution is very slow at best. If the oxide is heated with the alkali oxides or carbonate, the structure will be "opened out," and dilute acid will be sufficient. Samples containing other Group I or II elements will result in an ash that is soluble in 1:1 HCl, for example. Freshly precipitated iron oxides and hydroxides are all readily soluble in dilute acids. Although HCl is the best acid, the use of dilute nitric acid is common. HCl works better, due to the ligand properties not present with the nitrate anion. If the sample preparation requires that the sample be ashed, then the addition of an ashing aid, such as sodium, magnesium, or calcium carbonate, makes the dissolution step much easier, and this approach is compatible with Pt crucibles, but is not recommended for alumina, quartz, or porcelain crucibles. Many preparations of Fe-containing samples use either lithium carbonate, lithium borate (meta or tetra), or sodium carbonate fusions, and the ashing step is skipped altogether. If the sample contains a lot of organic or combustible material, then a programmable muffle furnace, where the temperature is slowly ramped to 1000 °C (in the case of sodium carbonate as the ashing aid), is prudent.

Ores — The sample should be pulverized to pass an 80 to 100 mesh sieve. Sulfides and ores containing organic matter can be mixed with sodium carbonate and ashed in a Pt crucible at 500 °C for 1 hour, prior to bringing to a temperature of 1000 °C as described below. If the ore has already been strongly ignited, it can go directly to 1000 °C with sodium carbonate (0.25 to 1.0 gram sample + 20 grams sodium carbonate, fused at 1000 °C for 10 to 30 minutes in a Pt crucible, followed by dissolution with 1:1 HCl). When a sodium carbonate fusion is performed, HF is not needed, unless silica content is high (silicates formed from the carbonate fusion are soluble, but do not heat the fuseate to dissolve, since this will bring about the formation of polysilicic acid, which is a gelatinous precipitate). This approach is applicable to sulfides, ores containing organic matter, oxides, red and brown hematites, magnetic iron ore, spathose iron ore, roasted pyrites, and iron ore briquettes.

Organic Matrices

Ashing of organic materials, foodstuffs, plant, and blood and sewage sludge, as a preliminary decomposition step, is not suggested for samples containing Fe, unless an ashing aid such as sodium carbonate is used and the crucible material is Pt. If magnesium nitrate is used as the ashing aid, then porcelain or quartz crucibles are acceptable. Conventional dry ashing procedures will result in the formation of a refractory form of iron oxide, unless there is a significant amount of one or more elements from Groups I, II, or III. If ashing is used, it is suggested to keep the temperature low (400 to 450 °C max) and to use an ashing aid, such as high-purity sodium carbonate or magnesium nitrate. If the sample is high in silica, subsequent fusion of the ash with sodium carbonate is suggested, or heating of the ash to fumes with sulfuric and hydrofluoric acids.

For more on ashing, please see the following paper: Ashing

The acid digestion of organic samples containing Fe is also very common. Fe does not offer any unusual challenges over the other elements and, with the exception of a higher-than-average contamination risk, is very well behaved. There are numerous oxidative acid digestion procedures available that are tailored to the matrix.

For more on acid digestion, please see the following paper: Acid Digestions of Organic Samples